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Early definitions of acids (sour taste, turning litmus red) were inadequate because they described properties, not structure. In 1923, Johannes Bronsted and Thomas Lowry independently proposed a definition based on proton transfer:
Every acid-base reaction is therefore a proton transfer from the acid to the base:
HA + B -> A- + BH+
The acid HA loses a proton to become A-. The base B gains a proton to become BH+. Notice that the reaction is reversible in principle - the products can themselves act as a (weaker) acid and base.
When hydrogen chloride dissolves in water the following equilibrium is set up (for HCl it lies far to the right):
HCl(aq) + H2O(l) -> H3O+(aq) + Cl-(aq)
H3O+ is the species that gives acidic solutions their properties. At A-Level we often simplify H3O+(aq) to H+(aq) for clarity, but you should recognise that the bare proton does not exist free in water; it is always solvated.
When an acid donates a proton, the species left behind can accept a proton back. This species is called the conjugate base of the acid. Likewise, when a base accepts a proton, the resulting species can donate it back - this is the conjugate acid of the base.
In HCl + H2O -> H3O+ + Cl-:
| Role | Species | Conjugate |
|---|---|---|
| Acid 1 | HCl | Cl- (conjugate base of HCl) |
| Base 1 | H2O | H3O+ (conjugate acid of H2O) |
Two conjugate acid-base pairs appear in every Bronsted-Lowry reaction. They differ by a single proton:
graph LR
A[HCl<br/>acid 1] -->|loses H+| B[Cl-<br/>base 1]
C[H2O<br/>base 2] -->|gains H+| D[H3O+<br/>acid 2]
The pairs are HCl / Cl- and H2O / H3O+. Each pair differs by exactly one H+.
NH3(aq) + H2O(l) <=> NH4+(aq) + OH-(aq)
Ammonia is a weak base, so the equilibrium lies to the left - the solution contains mostly NH3 with a small amount of NH4+ and OH-.
Notice that in the HCl example water behaves as a base (accepts a proton) but in the NH3 example water behaves as an acid (donates a proton). A species that can act as either a Bronsted-Lowry acid or a Bronsted-Lowry base is said to be amphoteric.
Water is the archetypal amphoteric solvent:
This dual behaviour underlies the self-ionisation of water (lesson 2, Kw).
Some acids can release more than one proton per molecule. The basicity of an acid is the number of replaceable hydrogen atoms (protons it can donate).
| Acid | Formula | Basicity | Ionisation steps |
|---|---|---|---|
| Hydrochloric | HCl | Monobasic (monoprotic) | HCl -> H+ + Cl- |
| Nitric | HNO3 | Monobasic | HNO3 -> H+ + NO3- |
| Ethanoic | CH3COOH | Monobasic | CH3COOH <=> H+ + CH3COO- |
| Sulfuric | H2SO4 | Dibasic (diprotic) | H2SO4 -> H+ + HSO4-; HSO4- <=> H+ + SO4^2- |
| Carbonic | H2CO3 | Dibasic | H2CO3 <=> H+ + HCO3-; HCO3- <=> H+ + CO3^2- |
| Phosphoric | H3PO4 | Tribasic (triprotic) | three successive ionisations |
Exam tip: not every H atom in an acid molecule is acidic. In ethanoic acid CH3COOH only the COOH hydrogen is acidic; the three CH3 hydrogens are not donated. The basicity is therefore 1, not 4. Acidic hydrogens are those attached to strongly electronegative atoms (O, halogens) where the O-H (or similar) bond is polar enough to release H+ in water.
All the familiar acid reactions can be seen as proton transfer from the acid to a base:
Mg(s) + 2HCl(aq) -> MgCl2(aq) + H2(g)
Here the metal is not a Bronsted-Lowry base; this is a redox reaction rather than a pure acid-base reaction. But the acid still behaves as a proton donor at the metal surface.
HCl(aq) + NaOH(aq) -> NaCl(aq) + H2O(l) Net ionic: H+(aq) + OH-(aq) -> H2O(l)
OH- accepts a proton to form water. This is the archetypal neutralisation.
2HCl(aq) + Na2CO3(aq) -> 2NaCl(aq) + H2O(l) + CO2(g) HCl(aq) + NaHCO3(aq) -> NaCl(aq) + H2O(l) + CO2(g)
The carbonate or hydrogencarbonate ion is the base. It accepts protons, the resulting H2CO3 decomposes rapidly to H2O and CO2, which is why effervescence is observed.
HCl(aq) + NH3(aq) -> NH4Cl(aq) Net ionic: H+(aq) + NH3(aq) -> NH4+(aq)
NH3 is the base.
In every case one species donates a proton and another accepts it. Identifying the pairs is the first step to writing any acid-base equation correctly.
Identify the conjugate acid-base pairs in:
HSO4-(aq) + H2O(l) <=> SO4^2-(aq) + H3O+(aq)
Solution:
Pairs: HSO4- / SO4^2- and H2O / H3O+.
Note that HSO4- itself is also the conjugate base of H2SO4. The same species can be an acid in one reaction and a base in another - it depends on the partner.
Bronsted-Lowry theory defines acid-base chemistry in terms of proton transfer: acids donate H+, bases accept H+. Each reaction contains two conjugate acid-base pairs, each differing by one H+. Water is amphoteric, acting as either acid or base depending on the partner. Basicity (mono-, di-, tri-) counts only the protons an acid can actually donate. These ideas are the foundation for quantitative pH, Ka and buffer calculations in the rest of this module.
Identify the acid, base and conjugates in:
NH4+(aq) + H2O(l) <=> H3O+(aq) + NH3(aq)
Solution: NH4+ donates a proton -> acid; H2O accepts a proton -> base. Pairs: NH4+ / NH3 and H2O / H3O+. Notice this is the opposite of the NH3 + H2O example earlier: here the ammonium ion acts as the acid because the proton is transferred from NH4+ to water.
Write the Bronsted-Lowry equation for the first step of the reaction between carbonate ion and ethanoic acid.
CH3COOH(aq) + CO3^2-(aq) <=> CH3COO-(aq) + HCO3-(aq)
Here CH3COOH is acid 1 (donates H+) and CO3^2- is base 2 (accepts H+). The conjugate pairs are CH3COOH / CH3COO- and HCO3- / CO3^2-. A second proton transfer can then produce H2CO3, which decomposes to CO2 and H2O - hence the familiar effervescence.
Before Bronsted and Lowry, Svante Arrhenius defined acids as species that dissociate in water to give H+ ions, and bases as species that dissociate to give OH-. This works for HCl, HNO3 and NaOH but fails to explain why NH3 (with no OH- in its formula) is a base, or why reactions can occur in non-aqueous solvents. The Bronsted-Lowry extension replaced "gives H+/OH- in water" with "donates/accepts protons" - a broader definition that works in any solvent and covers ammonia as a base. An even more general definition was later proposed by G. N. Lewis (acids as electron-pair acceptors), but at A-Level the Bronsted-Lowry view is usually sufficient.
In the next lesson we will meet the ionic product of water Kw, which turns the qualitative idea of water self-ionising into a quantitative equation we can use to calculate pH.